Chemical Compounds

CHEMTUTOR

IONIC AND COVALENT BONDS

A bond is an attachment among atoms. Atoms may be held together for any of several reasons, but all bonds have to do with the electrons, particularly the outside electrons, of atoms. There are bonds that occur due to sharing electrons. There are bonds that occur due to a full electrical charge difference attraction. There are bonds that come about from partial charges or the position or shape of electrons about an atom. But all bonds have to do with electrons. Since chemistry is the study of elements, compounds, and how they change, it might be said that chemistry is the study of electrons. If we study the changes brought about by moving protons or neutrons, we would be studying nuclear physics. In chemical reactions the elements do not change from one element to another, but are only rearranged in their attachments.

A compound is a group of atoms with an exact number and type of atoms in it arranged in a specific way. Every bit of that material is exactly the same. Exactly the same elements in exactly the same proportions are in every bit of the compound. Water is an example of a compound. One oxygen atom and two hydrogen atoms make up water. Each hydrogen atom is attached to an oxygen atom by a bond. Any other arrangement is not water. If any other elements are attached, it is not water. H₂O is the formula for that compound. This formula indicates that there are two hydrogen atoms and one oxygen atom in the compound. H₂S is hydrogen sulfide. Hydrogen sulfide does not have the same types of atoms as water. It is a different compound. H₂O₂ is the formula for hydrogen peroxide. It might have the right elements in it to be water, but it does not have them in the right proportion. It is still not water. The word formula is also used to mean the smallest bit of any compound. A molecule is a single formula of a compound joined by covalent bonds. The Law of Constant Proportions states that a given compound always contains the same proportion by weight of the same elements.

IONIC BONDS

Some atoms, such as metals tend to lose electrons to make the outside ring or rings of electrons more stable and other atoms tend to gain electrons to complete the outside ring. An ion is a charged particle. Electrons are negative. The negative charge of the electrons can be offset by the positive charge of the protons, but the number of protons does not change in a chemical reaction. When an atom loses electrons it becomes a positive ion because the number of protons exceeds the number of electrons. Non-metal ions and most of the polyatomic ions have a negative charge. The non-metal ions tend to gain electrons to fill out the outer shell. When the number of electrons exceeds the number of protons, the ion is negative. The attraction between a positive ion and a negative ion is an ionic bond. Any positive ion will bond with any negative ion. They are not fussy. An ionic compound is a group of atoms attached by an ionic bond that is a major unifying portion of the compound. A positive ion, whether it is a single atom or a group of atoms all with the same charge, is called a cation, pronounced as if a cat were an ion. A negative ion is called an anion, pronounced as if Ann were an ion. The name of an ionic compound is the name of the positive ion (cation) first and the negative (anion) ion second.

The valence of an atom is the likely charge it will take on as an ion. The names of the ions of metal elements with only one valence, such as the Group 1 or Group 2 elements, are the same as the names of the elements. A sodium atom that has lost the electron in its outside shell is a positive ion, sodium ion, Na⁺. A magnesium atom that has lost the two electrons in its outside shell is a plus two (double positive) magnesium ion, Mg⁺⁺.The names of the ions of nonmetal elements (anions) develop an -ide on the end of the name of the element. Nonmetal atoms tend to GAIN electrons, so when a nonmetal atom collects an extra electron, it will become a negative ion. For instance, fluorine ion is fluoride, F ^-; oxygen ion is oxide, O ⁼, a double negative ion because it gains TWO electrons; and iodine ion is iodide, I ^-. There are a number of elements, usually transition elements that have more than one valence and have a name for each ion, for instance, ferric ion is an iron ion with a positive three charge. Ferrous ion is an iron ion with a charge of plus two. There are a number of common groups of atoms that have a charge for the whole group. Such a group is called a polyatomic ion or radical. Chemtutor suggests it is best to learn by rote the list of polyatomic ions with their names, formulas and charges and the elements with more than one valence the same way.

SOME ATOMS WITH MULTIPLE VALENCES.

NOTE THERE ARE TWO COMMON NAMES FOR THE IONS. YOU SHOULD KNOW BOTH THE STOCK SYSTEM AND THE OLD SYSTEM NAMES.

ION	STOCK SYSTEM	OLD SYSTEM	ION	STOCK SYSTEM	OLD SYSTEM
Fe^2+	iron II	ferrous	Fe^3+	iron III	ferric
Cu^+	copper I	cuprous	Cu^2+	copper II	cupric
Au^+	gold I	aurous	Au^3+	gold III	auric
Sn^2+	tin II	stannous	Sn^4+	tin IV	stannic
Pb^2+	lead II	plumbous	Pb^4+	lead IV	plumbic
Hg^+	mercury I	mercurous	Hg^2+	mercury II	mercuric
Cr^2+	chromium II	chromous	Cr^3+	chromium III	chromic
Mn^2+	manganese II	manganous	Mn^3+	manganese III	manganic

The ions by the Stock system are pronounced, “copper one”, “copper two”, etc. Notice that the two most likely ions of an atom that has multiple valences have suffixes in the old system to identify them. The smallest of the two charges gets the “-ous” suffix, and the largest of the two charges has the “-ic” suffix. This leads to the amusing possibility of Saint Nickelous coming down your chimney. (Boo! Hiss!)

SOME ATOMS WITH ONLY ONE COMMON VALENCE:

• ALL GROUP 1 ELEMENTS ARE +1
• ALL GROUP 2 ELEMENTS ARE +2
• ALL GROUP 7 (HALOGEN) ELEMENTS ARE -1 WHEN IONIC
• Oxygen and sulfur (GROUP 6) are -2 when ionic
• Hydrogen is usually +1
• Al³⁺, Zn²⁺, and Ag⁺

RADICALS OR POLYATOMIC IONS

The following radicals or polyatomic ions are groups of atoms of more than one kind of element attached by covalent bonds. They do not often come apart in ionic reactions. The charge on the radical is for the whole group of atoms as a unit. These are common radicals you should learn WITH THEIR CHARGE AND NAME. HELP! THINGS LIKE THIS ARE WHY CHEMTUTOR WAS WRITTEN BAREFOOT IN HTML (WITHOUT AN HTML EDITOR). I JUST CANNOT SEEM TO GET THESE POLYATOMIC IONS LOOKING RIGHT WITH THEIR SUBSCRIPTS AND SUPERSCRIPTS. It seems like the more I play with them the messier they get. Also, the source code here has some considerable differences from the web. dmw TRY TO ERASE THE FORMATING WITH THE ERASE BUTTON ON THE EDITOR AND THEN USE THE SUPERSCRIPT AND SUBSCRIPT BUTTONS. dsl

• (NH₄)^₊ AMMONIUM - Do not confuse with NH₃, AMMONIA GAS)

• (NO₃)^- NITRATE (Do not confuse with NITRIDE (N^3-) or NITRITE)

• (NO₂)^- NITRITE (Do not confuse with (N^3-) or NITRATE)

• (C₂H₃O₂)^- ACETATE (NOTE - This is not the only way this may be written.)

• (ClO₃)^- CHLORATE (Do not confuse with CHLORIDE (Cl^-) or CHLORITE)

• (ClO₂)^- CHLORITE (Do not confuse with CHLORIDE (Cl^- ) or CHLORATE)

• (SO3)^2- SULFITE (Do not confuse with (S^2-) or SULFATE)

• (SO4)^2- SULFATE (Do not confuse with SULFIDE (S^2-) or SULFITE)

• (HSO3)^- BISULFITE (or HYDROGEN SULFITE)

• (PO4)^3- PHOSPHATE (Do not confuse with P^3-, PHOSPHIDE)

• (HCO3)^- BICARBONATE (or HYDROGEN CARBONATE)

• (CO3)^2- CARBONATE

• (HPO4)^2- HYDROGEN PHOSPHATE

• (H2PO4)^- DIHYDROGEN PHOSPHATE

• (OH)^- HYDROXIDE

• (CrO4)^2- CHROMATE

• (Cr2O7)^2-DICHROMATE

• (BO3)^3- BORATE

• (AsO4)^3- ARSENATE

• (C2O4)^2- OXALATE

• (ClO4)^- PERCHLORATE

• (CN)^- CYANIDE

• (MnO4)^- PERMANGANATE

ACIDS OF SOME COMMON POLYATOMIC IONS.

These are written here with the parentheses around the polyatomic ions to show their origin. Usually these compounds are written without the parentheses, such as HNO₃ or H₂SO₄. Note that the acids of polyatomic ions with a single negative charge only have one hydrogen. Polyatomic ions with two negative charges have two hydrogens.

• H(OH) WATER (!)
• H(NO3) NITRIC ACID
• H(NO2) NITROUS ACID .
• H(C2H3O2) ACETIC ACID
• H2(CO3) CARBONIC ACID
• H2(SO3) SULFUROUS ACID
• H2(SO4) SULFURIC ACID
• H3(PO4) PHOSPHORIC ACID
• H2(CrO4) CHROMIC ACID
• H3(BO3) BORIC ACID
• H2(C2O4) OXALIC ACID

WRITING IONIC COMPOUND FORMULAS

In the lists above, the radicals and compounds have a small number after and below an element if there is more than one of that type of that atom. For instance, ammonium ions have one nitrogen atom and four hydrogen atoms in them. Sulfuric acid has two hydrogens, one sulfur, and four oxygens.

Knowing the ions is the best way to identify ionic compounds and to predict how materials would join. People who do not know of the ammonium ion and the nitrate ion would have a difficult time seeing that NH₄NO₃ is ammonium nitrate. Chemtutor very highly recommends that you know all the above ions, complete with the valence or charge. One of the best ways to learn the ions is to write ionic compounds.

Let’s consider what happens in an ionic bond using electron configuration, the octet rule, and some creative visualization. A sodium atom has eleven electrons around it. The first shell has two electrons in an s subshell. The second shell is also full with eight electrons in an s and a p subshell. The outer shell has one lonely electron, as do the other elements in Group 1. This outside electron can be detached from the sodium atom, leaving a sodium ion with a single positive charge and an electron. A chlorine atom has seventeen electrons. Two are in the first shell, eight are in the second shell, and seven are in the outside shell. The outside shell is lacking one electron to make a full shell, as are all the elements of Group 7. When the chlorine atom collects another electron, the atom becomes a negative ion. The positive sodium ion missing an electron is attracted to the negative chloride ion with an extra electron. The symbol for a single unattached electron is e^-.

½Cl2 + Na Cl + e^- + Na⁺ Cl^- + Na⁺ Na⁺Cl^- NaCl

Any compound should have a net zero charge. The single positive charge of the sodium ion cancels the single negative charge of the chloride ion. The same idea would be for an ionic compound made of ions of plus and minus two or plus and minus three, such as magnesium sulfate or aluminum phosphate

Mg²⁺ + (SO₄)^2- Mg²⁺(SO₄)^2- Mg(SO₄) or MgSO4

Al³⁺ + (PO4)^3- Al³⁺(PO₄)^3- Al(PO₄) or AlPO₄

But what happens if the amount of charge does not match? Aluminum bromide has a cation that is triple positive and an anion that is single negative. The compound must be written with one aluminum and three bromide ions. AlBr₃. Calcium phosphate has a double positive cation and a triple negative anion. If you like to think of it this way, the number of the charges must be switched to the other ion. Ca₃(PO₄)₂. Note that there must be two phosphates in each calcium phosphate, so the parentheses must be included in the formula to indicate that. Each calcium phosphate formula (Ionic compounds do not make molecules.) has three calcium atoms, two phosphphorus atoms, and eight oxygen atoms.

There are a small number of ionic compounds that do not fit into the system for one reason or other. A good example of this is magnetite, an ore of iron, Fe₃O₄. The calculated charge on each iron atom would be +8/3, not a likely actual charge. The deviance from the system in the case of magnetite could be accounted for by a mixture of the common ferric and ferrous ions.

BINARY COVALENT COMPOUNDS

The word binary means that there are two types of atom in a compound. Covalent compounds are groups of atoms joined by covalent bonds. Binary covalent compounds are some of the very smallest compounds attached by covalent bonds. A covalent bond is the result of the sharing of a pair of electrons between two atoms. The chlorine molecule is a good example of the bond, even if it has only one type of atom. Chlorine gas, Cl₂, has two chlorine atoms, each of which has seven electrons in the outside ring. Each atom contributes an electron to an electron pair that make the covalent bond. Each atom shares the pair of electrons. In the case of chlorine gas, the two elements in the bond have exactly the same pull on the electron pair, so the electrons are exactly evenly shared. The covalent bond can be represented by a pair of dots between the atoms, Cl:Cl, or a line between them, Cl-Cl. Sharing the pair of electrons makes each chlorine atom feel as if it has a completed outer shell of eight electrons. The covalent bond is much harder to break than an ionic bond. The ionic bonds of soluble ionic compounds come apart in water, but covalent bonds do not usually come apart in water. Covalent bonds make real molecules, groups of atoms that are genuinely attached to each other. Binary covalent compounds have two types of atom in them, usually non-metal atoms. Covalent bonds can come in double (sharing of two pairs of electrons) and triple (three pairs of electrons) bonds.

FORMULA	COMMON NAME	SYSTEM NAME
N2O	nitrous oxide	dinitrogen monoxide
NO	nitric oxide	nitrogen monoxide
N2O3	nitrous anhydride	dinitrogen trioxide
NO2	nitrogen dioxide	nitrogen dioxide
N2O4	nitrogen tetroxide	dinitrogen tetroxide
N2O5	nitric anhydride	dinitrogen pentoxide
NO3	nitrogen trioxide	nitrogen trioxide

With the compounds of nitrogen and oxygen to use as examples, we see that there are often more ways for any two elements to combine with each other by covalent bonds than by ionic bonds. Many of the frequently seen compounds already have names that have been in use for a long time. These names, called common names, may or may not have anything to do with the makeup of the material, but more of the common names of covalent compounds are used than of the ionic compounds.

*FGP	number	*FGP	number	*FGP	number	*FGP	number
mono-	one	di-	two	tri-	three	tetra-	four
penta-	five	hexa-	six	hepta-	seven	octa-	eight
nona-	nine	deca-	ten	undeca-	eleven	dodeca-	twelve

The system names include numbers that indicate how many of each type of atom are in a covalent molecule. The Fake Greek Prefixes (FGP’s above in the chart) are used to indicate the number. It would be wise of you to know the FGP’s.

In saying or writing the name of a binary covalent the FGP of the first element is said, then the name of the first element is said, then the FGP of the second element is said, and the name of the second element is said, usually with the ending “-ide” on it. The only notable exception for the rule is if the first mentioned element only has one atom in the molecule, in which case the “mono-“ prefix is omitted. CO is carbon monoxide. CO₂ is carbon dioxide. In both cases there is only one carbon in the molecule, and the “mono-“ prefix is not mentioned. For oxygen the last vowel of the FGP is omitted, as in the oxides of nitrogen in the above table.

COMMON NAMES OF BINARY COVALENT COMPOUNDS YOU SHOULD KNOW

• H₂O water
• N₂H₄ hydrazine
• CH₄ methane
• C₂H₂ acetylene
• NH₃ ammonia

CHECKLIST OF KNOWLEDGE FOR WRITING COMPOUNDS

Here’s a checklist of the things you need to know to be able to correctly write the formulas for materials.

• NAMES AND SYMBOLS OF THE ELEMENTS
• NAMES AND SYMBOLS OF DIATOMIC GASES
• NAMES, SYMBOLS, AND VALENCES OF THE ELEMENTS IN GROUPS 1, 2, 7, and 8
• NAMES, SYMBOLS, AND VALENCES OF METALS WITH ONE COMMON VALENCE
• NAMES AND VALENCES OF METALS WITH MORE THAN ONE COMMON VALENCE
• NAMES, FORMULAS, AND CHARGES OF COMMON POLYATOMIC IONS
• NAMES AND FORMULAS OF COMMON ACIDS
• HOW TO TELL THE DIFFERENCE BETWEEN COVALENT AND IONIC COMPOUNDS
• HOW TO WRITE THE FORMULA OF IONIC COMPOUNDS
• LIST OF FAKE GREEK PREFIXES UP TO TWELVE
• HOW TO WRITE THE FORMULA OF BINARY COVALENT COMPOUNDS
• COMMON NAMES OF SOME BINARY COVALENT COMPOUNDS

MORE ON BONDS, SHAPES, AND OTHER FORCES

THE CONTINUUM BETWEEN IONIC AND COVALENT BONDS

In an attempt to simplify, some books may seem to suggest that covalent and ionic bonds are two separate and completely different types of attachment. A covalent bond is a shared pair of electrons. The bond between the two atoms of any diatomic gas, such as chlorine gas, Cl₂, is certainly equally shared. The two chlorine atoms have exactly the same pull on the pair of electrons, so the bond must be exactly equally shared. In cesium fluoride the cesium atom certainly donates an electron and the fluoride atom certainly craves an electron. Both the cesium ion (Cs⁺) and the fluoride (F^-) ion can exist in solution independently of the other. The bond between a cesium and a fluoride ion to make cesium fluoride (CsF) would be clearly ionic because the difference in electronegativities (ΔEN) is so large.

The amount of pull on an atom has on a shared pair of electrons, called electronegativity, is what determines the type of bond between atoms. Considering the Periodic Table without the inert gases, electronegativity is greatest in the upper right of the Periodic Table and lowest at the bottom left. The bond in francium fluoride should be the most ionic. Some texts refer to a bond that is between covalent and ionic called a polar covalent bond. There is a range of bond between purely ionic and purely covalent that depends upon the electronegativity of the atoms around that bond. If there is a large difference in electronegativity, the bond has more ionic character. If the electronegativity of the atoms is more similar, the bond has more covalent character.

LEWIS STRUCTURES

Lewis structures are an opportunity to better visualize the valence electrons of elements. In the Lewis model, an element symbol is inside the valence electrons of the s and p subshells of the outer ring. It is not very convenient to show the Lewis structures of the Transition Elements, the Lanthanides, or Actinides. The inert gases are shown having the element symbol inside four groups of two electrons symbolized as dots. Two dots are above the symbol, two below, two on the right, and two on the left. The inert gases have a full shell of valence electrons, so all eight valence electrons appear. Halogens have one of the dots missing. It does not matter on which side of the symbol the dot is missing. Group 1 elements and hydrogen are shown with a single electron in the outer shell. Group 2 elements are shown with two electrons in the outer shell, but those electrons are not on the same side. Group 3 elements have three dots representing electrons, but the electrons are spread around to one per position, as in Group 2 elements. Group 4 elements, carbon, silicon, etc. are shown as having four electrons around the symbol, each in a different position.

Group 5 elements, nitrogen, phosphorus, etc. have five electrons in the outer shell. In only one position are there two electrons. So Group 5 elements such as nitrogen can either accept three electrons to become a triple negative ion or join in a covalent bond with three other items. When all three of the unpaired electrons are involved with a covalent bond, there is yet another pair of electrons in the outside shell of Group 5 elements.

Group 6 elements, oxygen, sulfur, etc., have six electrons around the symbol, again without any concern to position except that there are two electrons in two positions and one electron alone in the other two positions. Group 7 elements have all of the eight outside electrons spaces filled except for one. The Lewis structure of a Group 7 element will have two dots in all four places around the element symbol except for one.

Let's start with two atoms of the same type sharing a pair of electrons. Chlorine atoms have seven electrons each and would be a lot more stable with eight electrons in the outer shell. Single chlorine atoms just do not exist because they get together in pairs to share a pair of electrons. The shared pair of electrons make a bond between the atoms. In Lewis structures, the outside electrons are shown with dots and covalent bonds are shown by bars.

This covalent bond between chlorine is one of the most covalent bonds known. Why? A covalent bond is the sharing of a pair of electrons. The two atoms on ether side of the bond are exactly the same, so the amount of "pull" of each atom on the electrons is the same, and the electrons are shared equally.

Next let's consider a molecule in which the atoms bonded are not the same, but the bonds are balanced. Methane, CH₄, is such a molecule. If there were just a carbon and a single hydrogen, the bond between them would not be perfectly covalent. In the CH₄ molecule, the four hydrogen atoms exactly balance each other out. The Lewis structure of methane does not have any electrons left over. The carbon began with four electrons and each hydrogen began with two electrons. Only the bars representing the shared pairs of electrons remain. The carbon now shares four pairs of electrons, so this satisfies the carbon's need for eight electrons in the outside shell. Each hydrogen has a single shared pair in the outside shell, but the outside shell of the hydrogen only has two electrons, so the hydrogen has a full outer shell also.

Carbons and hydrogens are nice and easy to write in Lewis structures, because each carbon must have four attachments to it and each hydrogen atom must have one and only one attachment to it. When the bonds around a carbon atom go to four different atoms, the shape of the bonds around that carbon is roughly tetrahedral, depending upon what the materials are around the carbon. Carbons are also able to have more than one bond between the same two. Consider the series ethane (C₂H₆), ethene (C₂H₄),(common name is ethylene), and ethyne (C₂H₂), (common name is acetylene).

H₃ – C – C – H₃ ethane H₂ – C = C – H₂ ethylene H – C ≡ C – H acetylene

In writing the Lewis structure of compounds, the bars representing bonds are preferred to the dots representing individual electrons.

The double bars between the carbons in ethylene, C=C, represent a double bond between the two carbons, that is four shared electrons to make a stronger attachment between the two carbons. The triple bars between the carbons of acetylene represent a triple covalent bond between those two carbons, C≡C, three pairs of shared electrons between those carbons. Every carbon has four bonds to it showing a pair of electrons to make eight electrons (or four orbitals) in the outer shell. Each hydrogen atom has one and only one bond to it for two electrons in the outer shell that occupies the only orbital that hydrogen has. All of the outer shells are usually filled.

While we are doing this, notice that the Lewis structure of a molecule will show the shape of the molecule. All of the bonds in ethane are roughly the tetrahedral angle, so all of the hydrogen atoms are equivalent. This is true. The bonds in acetylene make it a linear molecule. The bonds in ethylene are somewhat trigonal around the carbons, and the carbons cannot twist around that bond as they can around a single bond, so that the molecule has a flat shape and the attachments to the carbons are not equivalent. This is also true. (You will see this in the study of organic chemistry. This type of difference between the positions of the hydrogen atoms is called cis - trans isomerism.)

The Lewis structure shows the shape of a molecule or polyatomic ion with the bonds to each atom drawn at 90 degrees (right, left, up, and down) from the atomic symbol and the non – bonded electrons as dots, usually in pairs, around the atomic symbol in the left, right, up, and down positions around the atom. We could set up a group of general guidelines for the drawing of Lewis structures for simple molecules or polyatomic ions.

Write all the atoms in the material in the form of the formula of the compound. CO₂ can be an example.

• Usually pick the atom with the lowest electronegativity (most distant from fluorine on the Periodic Table) to be the central atom or atoms. (In most organic compounds, carbon provides the main "skeleton" of the molecule.) The lowest electronegativity atom, the central atom, is usually written first in the compound. Carbon is the obvious candidate for the central atom.

• Arrange the other atoms around the inner core according the formula of the material using single bonds to hold the structure together. This is called the skeleton structure. The skeleton structure for carbon dioxide should be: O – C – O

• Count the Total Valence Electrons (TVE) of the molecule. This is done by adding up the electrons in the outside shell of each atom. This is easy for “main sequence” atoms, groups I, II, IIIA, IVA, VA, VIA, and VIIA. (See the Periodic Table as you do this.) Hydrogen and group I, or group 1, all have 1 electron in the outside shell. Group II or 2, starting with beryllium, have two electrons in the outer shell. Group IIIA or 13, starting with boron, have three electrons in the outer shell. Group IVA or 14, starting with carbon have four electrons in the outer shell. Group VA or 15, beginning with nitrogen, have five electrons in the outer shell. Group VIA or 16, beginning with oxygen, have six electrons in the outer shell. Group VIIA or 17, the halogens, have seven electrons in the outer shell. Notice that we usually don’t include the eight electrons in the outer shell of most inert gases because the noble elements do not usually make compounds. There are six valence electrons in each oxygen atom in CO₂ for a total of 12 and four electrons in the carbon atom for a grand total of 16 electrons in the CO₂ structure. TVE = 16 electrons. Carbon dioxide is a molecule and does not have a charge, but if you draw the Lewis structure of a polyatomic ion, you should add an electron for each negative charge and remove an electron from the TVE for each positive charge.

• Subtract the number of electrons in the bonds of the skeleton structure from the TVE and you will have the number of electrons you have to represent as dots around the atoms. For CO2, the math is:

TVE = 16 electrons

Electrons in bonds = - 4 electrons (two bonds)

Dots needed = 12 dots

• Distribute the dots (representing electrons) around the structure to the terminal atoms first. Hydrogen does not get any dots. It has all the electrons it can take with just the bond. All other atoms get a maximum of four orbitals, six dots if the atom has one bond to it, four dots if the atom has two bonds to it, two dots if the atom has three bonds to it, and no bonds if it has four bonds to it.

. . . .
: O – C – O : This is the proposed shape for the CO₂ molecule in the skeletal form.
. . . .

• The proposed shape above has some problems with it. There are too many electrons assigned to the oxygen atoms and not enough to the carbon. The way to express this idea is the formal charge. The formal charge is the number of electrons the atom brought to the structure minus the number of electrons shown in the proposed structure. The oxygen atoms both had six electrons in the valence shell because they are group VI A or group 16 atoms. They SHOW seven electrons in the proposed scheme, six dots and one electron from half the bond. 6 – 7 = - 1, so the formal charge of both the oxygen atoms is -1. The carbon atom brought four electrons, being from group IV A or 14. Carbon shows only two electrons, one from each of the bonds, so 4 – 2 = 2. The formal charge of the carbon is plus two. The difference in formal charge indicates that there is a problem, but it also shows a likely way to balance things out.

• If you have a structure where there are atoms around a bond that have opposing charges, the likely way to even out those charges is to take a pair of electrons from the negative atom and make it part of a multiple bond with the positive atom. Now the CO₂ molecule looks a lot better. We changed the single bond to a double bond on both sides of the carbon. Now the formal charge of all three atoms is zero (You check it yourself.), and there are four and only four orbitals around each atom. Each oxygen atom has two bonds and two unshared (lone) pairs of electrons for a total of four orbitals. The carbon has four bonds to it, four orbitals. This condition with the lowest number of formal charges and the right number of orbitals around each atom is the most stable and the most likely correct Lewis structure.

. . . .
: O = C = O :

This process of writing Lewis structures is very limited to small molecules. There are many exceptions to the process, for instance, there are some compounds in which one atom has only three orbitals around it. BF3, boron trifluoride is one in which the boron atom (central) is stuck with just three bonds to it. Some central atoms can have MORE than four orbitals around them. There is a phosphorus trichloride molecule (PCl₃) that has the same shape as ammonia, but there is also a phosphorus pentachloride molecule (PCl₅) that has five chlorine atoms attached to a central phosphorus. As you see, the scope of this tutorial goes only so far into the Lewis structure world.

With the warnings in mind, here are some general rules that will generally help you write Lewis structures.

• • HONC, pronounced “honk.” This is the way to remember that all Hydrogens have one and only one bond to them. Most Oxygens have two bonds to them. Most Nitrogens have three bonds to them, and most Carbons have four bonds to them.
  SO REMEMBER: HONC 1, 2, 3, 4.
• • Carbon is always a central atom, except in diatomic molecules like carbon monoxide.
• • Hydrogen is always a terminal atom with only one bond and no dots.
• • The lowest electronegativity atom (NOT the closest to fluorine) is usually the central atom.
• • The structures are usually pretty well balanced around the central atom.

The most likely - looking Lewis structure of a molecule is usually a good indicator of the actual shape of the molecule. We can tell that from the properties of the molecules. Rarely, but sometimes the best – looking Lewis structure is not the structure that predicts the properties of the material. In this case, the Lewis structure is wrong, and it probably makes some sense once the Lewis structure is written in the way that goes with the properties of the material.

SHAPES AROUND AN ATOM, VSEPR THEORY

There is no issue of shape around the Group 1 elements. There is only one attachment to them, so no angle is possible around them. But there are some molecular compounds with only two atoms, such as nitrogen monoxide, NO. The only feature of this molecule is the bond between the nitrogen atom and the oxygen atom. The small difference in electronegativity between the oxygen and the nitrogen give the molecule a small dipole, a small separation of charge, so a small amount of polarity. (Because there are an odd number of electrons in NO, this makes for an interesting Lewis structure. Try it.) Iodine fluoride, IF, is another diatomic compound that should have some polarity. Diatomic molecules like chlorine gas, Cl₂, have no electronegativity difference (ΔEN) from side to side of the bond, so they are completely balanced and completely non – polar.

Group 2 elements have two electrons in the outer shell. Many of the compounds of Group 2 elements are ionic compounds, not really making an angle in a molecule. Molecules made with Group 2 elements that have two attached items to the Group 2 element have a linear shape, because the two attached materials will try to move as far from each other as possible. A linear shape means that a straight line could be made through all three atoms with the central element in the center. The shape of carbon dioxide is linear with the carbon in the center. O = C = O

VSEPR stands for Valence Shell Electron Pair Repulsion. The idea is a disarmingly simple one. Electrons are all negatively charged, so they repel each other. If an atom has two electron groups around it, the electrons, and the atoms they are bonded to, are likely to be found as far as they can be from each other. “As far as they can get from each other,” and still remain attached to the central atom means that the angle around the central atom is 180 degrees, a straight line. Molecules with two electron groups attached to a central atom have a linear electron group shape and a linear molecular shape. Unless there is a large difference in electronegativity from one side to the other of a linear compound, there is no separation of charge and no polar character of the molecule.

Covalent compounds with boron are good examples of trigonal shaped molecules. The trigonal shape is a flat molecule with 120 degree angles between the attached atoms. Again using the example of a boron atom in the center, the attached elements move as far away from each other as they can, forming a trigonal shape, also called triangular, or trigonal planar to distinguish it from the trigonal pyramidal shape of compounds like ammonia. BF3, boron trifluoride, is an example of a molecule with a trigonal planar shape. Each fluorine atom is attached to the central boron atom. There are three bonds to the boron, so the electron group shape is trigonal planar around boron. The molecular shape is also trigonal planar in boron trifluoride because each electron group has a fluorine atom attached to it.

But, what if the central atom has two other atoms and a lone pair of electrons attached to it? Nitrogen oxychloride is an example of that. NOCl, is a molecule with nitrogen in the center (See how to write Lewis structures above.) and an oxygen and a chlorine atom attached to the central nitrogen. When we go through the skeleton structure and distribute the electron dots, we find that there is a double bond between the nitrogen and the oxygen and a lone pair (unshared pair) of electrons on the nitrogen in addition to the single bond from the nitrogen to the chlorine. There are three electron groups around the nitrogen, making the electron group shape more or less trigonal planar. But only two of those electron groups have an atom attached, so the molecular shape of nitrogen oxychloride is bent or angular. NOCl is not a balanced shape, so it is likely that there is some separation of charge within the molecule, making it a somewhat polar compound.

Group 4 elements are not in the center of a flat molecule when they have four equivalent attachments to them. As with two or three attachments, the attached items move as far as they can away from each other. In the case of a central atom with four things attached to it, the greatest angle between the attached items does not produce a flat molecule. If you were to cut off the vertical portion of a standard three-legged music stand so that it was the same length as the three legs, the angles among all four directions would be roughly equal. Try this with a gumdrop or a marshmallow. Stick four different colored toothpicks into the center at approximately the same angle. If you have done it right, the general shape of the device will be the same no matter which one of the toothpicks is up. This shape is called tetrahedral. The shape of a tetrahedron appears with the attached atoms at the points of the figure and each triangle among any three of them makes a flat plane. A tetrahedron is a type of regular pyramid with a triangular base. Carbon is a group four element. Organic and biochemical compounds have carbon as a “backbone,” so this tetrahedral shape is very important. Methane, CH₄, and carbon tetrachloride, CCl₄, are good examples of tetrahedral shape. If you draw the Lewis structures of these compounds, you will see that there are four bonds to the central carbon atom, but no other electrons on the central atom. They have four electron groups (single bonds) around the central atom, so they have a tetrahedral electron group shape. Each bond to the central carbon has an atom attached, so they have a tetrahedral molecular shape. In both compounds, the four atoms attached to carbon are the same, so there is no separation of charge. All four atoms have the same electron pull in balanced directions, so these compounds are non – polar. Can a central carbon make molecules with other shapes around the central atom? Yes, you remember carbon dioxide, where there are two double bonds around the carbon.

O = C = O Each double bond is an electron group, so there are only two electron groups around the carbon in carbon dioxide.

See the “acid carbons,” the ones with the ionizable hydrogen (in blue) on it. The shape of around the acid carbons is trigonal planar because it has a double bond to it and only three electron groups, but the shape around the other carbons is tetrahedral. In the Lewis structures the atoms are drawn at ninety degrees from each other, but the real shape around those carbons exists in three – space.

Group 5 elements, for instance nitrogen or phosphorus, will become triple negative as they add three electrons in ionic reactions, but this is rare. Nitrides and phosphides do not survive in the presence of water. Covalent bonds with these elements do survive in water. From the Lewis structure of these elements in the previous section, you know that Group 5 elements have the capability of joining with three covalent bonds, but they don’t make the trigonal shape because the UNSHARED PAIR OF ELECTRONS ACTS LIKE ANOTHER BONDED ATTACHMENT. The shape of the bonds and the lone pair of electrons around nitrogen and phosphorus is tetrahedral, just like the bonds around Group 4 elements. The molecular shape is trigonal pyramidal.

Group 6 elements, oxygen and sulfur, have six electrons in the valence shell. The compounds they make usually have two pairs of unshared electrons. Just as in Group 5 elements, these pairs of unshared electrons serve as other attached atoms for the electron shape of the molecule. Group 6 elements make tetrahedral electron shapes, but now there are only two attached atoms. The angle between the hydrogens in water is about 105 degrees. This peculiar shape is one of the things that makes water so special.

Group 7 elements have only one chance of attachment, so there is not usually any shape around these atoms.

BONDING FORCES IN WATER

The alchemists of old had several other objectives aside from making gold. The thought of a fluid material that could dissolve anything, the universal solvent, was another alchemical project. No alchemist would say, though, what material would hold such a fluid. Surprisingly, the closest thing we have to a universal solvent is water. Water is not only a common material, but the range of materials it dissolves is enormous. The guiding principle for predicting which materials dissolve in which solvent is that 'like dissolves like.' Fluids in which the atoms are attached with covalent bonds will dissolve covalent molecules. Fluids with a separation of charge in the bonds will dissolve ionic materials.

The bonds that hold hydrogen atoms to oxygen atoms are closer to covalent than ionic, but the bond does have a great deal of ionic character. Oxygen atoms are more electronegative than hydrogen atoms, so the electron pair is held closer to the oxygen atom. Another way to look at it is that only a very small number of water molecules are ionized at any one time. The ionization of water, H₂O → H⁺ + (OH)^- , into hydrogen ions (actually, hydronium ions) and hydroxide ions happens in only a very small number of the water molecules, but the effect is quite important as the reason for the existence of acids and bases. Materials of a mildly covalent nature, such as small alcohols and sugars, are soluble in water due to the mostly covalent nature of the bonds in water.

The shape of the water molecule is bent at about a 105 degree angle due to the electron structure of oxygen. The two pairs of electrons that force the attached hydrogens into something close to a tetrahedral angle give the water molecule an unbalanced shape like a boomerang, with oxygen at the angle and the hydrogen atoms at the ends. We can think of the molecule has having an ‘oxygen side’ and a ‘hydrogen side’. Since the oxygen atom pulls the electrons closer to it, the oxygen side of the molecule has a slight negative charge. Cations (positive ions) are attracted to the partial positive charge on the oxygen side of water molecules. Likewise, the hydrogen side of the molecule has a slight positive charge, attracting anions. Polar materials such as salts, materials that have a separation of charge, dissolve in water due to the charge separation of water. The origin of the separation is called a dipole moment and the molecule itself can be called a dipole.

Molecules or atoms that have no center of asymmetry are non-polar. Atoms such as the inert gases have no center of asymmetry. Molecules such as methane, CH₄, are likewise totally symmetrical. Very small forces, called London forces, can be developed within such materials by the momentary asymmetries of the material and induction forces on neighboring materials. These small forces account for the ability of non-polar particles to become liquids and solids. The larger the atom or molecule, the more potent the London forces, possibly due to the greater ability to separate charge within a larger particle. The larger the inert gas, the higher its melting point and boiling point. In alkanes, a series of non-polar hydrocarbon molecules, the larger the molecule, the higher the melting and boiling point.

There may be London forces in water molecules, but the enormous force of the dipole interaction completely hides the small London forces. The dipole forces within water are particularly strong for two additional reasons. Dipole forces that involve hydrogen atoms around a strongly electronegative material such as nitrogen, oxygen, fluorine, or chlorine are particularly strong due to the small size of the hydrogen atom compared to the size of the dipole force. Such dipoles have significantly stronger forces, and have been called hydrogen bonds. In water, this effect is even greater due to the small size of the oxygen atom, thus the whole water molecule. In a water molecule hydrogen bonding is a large intermolecular force in a small volume on a small mass that makes it particularly noticeable.

Compare methane, CH4, to water. They are similar in size and mass, but methane is non-polar and water is very highly polar due to the hydrogen bonding. The melting point for methane is -184 °C (89 K) and for water is 0 °C (273 K). The boiling point for methane is -161.5 °C (111.7 K) compared to water at 100 °C (373.2 K). The temperature range over which methane is a liquid is less than a quarter the range for water. Most of these differences are accountable from the hydrogen bonding of water.

The properties of water come directly from the molecular shape of it and the forces it has on it from that shape. Water is cohesive. It balls up with itself in zero gravity or on a non – polar surface like waxed paper. The surface tension of water is another product of the cohesive forces, mainly hydrogen bonding. Water is adhesive, that is, it clings to other things. It wets cotton or paper, it wets glass or ceramic, and it dissolves many compounds, to include polar compounds.

Water is a very important material for living things because:

• • It has a high heat capacity, or specific heat; water absorbs or releases large amounts of heat with small changes in temperature.
• • It has a large range of temperature in which it is a liquid.
• • It has a high heat of vaporization; it takes a lot of heat to change liquid water into steam.
• • It is one of the best solvents, particularly for ionic materials.
• • Water forms hydration layers around large charged particles like proteins and nucleic acids that make the functions of the macromolecules possible.
• • It serves as the body’s major transport medium.
• • Water is an important part of hydrolysis and dehydration synthesis reactions.
• • Water forms a resilient cushion around certain body organs.

There are three main types of bonding forces, forces that make compounds. Ionic bonding is just the attraction of a positive ion for a negative ion. Sodium chloride is a compound that is made of sodium ions, having lost an electron, with a positive charge, and negative chloride ions, negative because they attract another electron to fill the valence shell. Covalent bonds come about by a bonded pair of atoms sharing a pair (or more pairs) of electrons. Covalent bonds are usually stronger than ionic bonds. Ionic bonds can separate in water solution. Polar covalent bonds, such as the bonds between the hydrogen and oxygen atoms of water, happen when two atoms sharing a pair of electrons have a large difference in electronegativity.

Three main types of intermolecular forces, hydrogen bonding, dipole interactions, and dispersion forces, are forces that do not make compounds, but attract or repel on an atomic level. The name London forces (from Fritz London) is sometimes used for the small dipole interactions and even smaller dispersion forces. Dispersion forces are caused by the momentary unbalance of electrons around an atom. They are called “dispersion” forces for the uneven dispersion of electrons. Even noble gases can have these forces. In fact dispersion forces are the only forces that pull noble gases together. In atoms or small molecules, dispersion forces are very small. The melting and boiling points of noble gases are very low because it takes very little energy to overcome the dispersion forces. In macromolecules, though, the dispersion forces can develop to be much larger. In proteins and nucleic acids, dispersion forces rival the magnitude of the dipole forces and even hydrogen bonding.

Dipole forces, or dipole – dipole interactions are the forces from polar molecules pulling together by the difference in charge from one side of a molecule to another. Iodine fluoride, IF, is likely to have a small positive charge near the iodine and a small negative charge near the fluorine, because fluorine is by far the most electronegative. The IF molecules have a tendency to arrange themselves with the positive end of one molecule near the negative end of another molecule. The dipole forces of water are fairly large due to the highly polar nature of the water molecule.

In water, the most powerful intermolecular force is hydrogen bonding. Hydrogen bonding is the tendency of hydrogen atoms attached to highly electronegative atoms like fluorine, chlorine, or oxygen to seek other highly electronegative atoms in other molecules. The forces can make liquids viscous and cohesive. Water owes its cohesive properties mostly to hydrogen bonding. But hydrogen bonding is even more important in macromolecules. The secondary, tertiary, and quaternary structures of macromolecules are due in large part to hydrogen bonding. The association of opposing nucleotides in nucleic acids is due to hydrogen bonding.

COMPOUND WORKSHEET

Write chemical formula as requested. Show subscript numbers where needed. Show valences for all ions.

1. hydrochloric acid _________________ 2. sodium chloride ________________ 1. HCl 2. NaCl

3. uranium hexafluoride _____________ 4. strontium nitrate ________________ 3. UF6 4. Sr(NO3)2

5. calcium chloride _________________ 6. acetic acid ___________________ 5.CaCl2 6.HC2H3O2

7. phosphoric acid __________________ 8. ammonia ______________________ 7. H3PO4 8. NH3

9. chlorine ______________________ 10. lithium sulfate ___________________ 9. Cl2 10. Li2SO4

11. potassium chromate ____________ 12. calcium hydroxide ____________ 11.K2CrO4 12.Ca(OH)2

13. aluminum foil _________________ 14. ammonium sulfate ______________ 13.Al 14.(NH4)2SO4

15. sulfuric acid __________________ 16. ammonium iodide ______________ 15. H2SO4 16. NH4I

17. acetylene _____________________ 18. rubidium nitrite _______________ 17. C2H2 18. RbNO2

19. lead II sulfite __________________ 20. copper I sulfide ________________ 19. PbSO3 20.Cu2S

21. aluminum oxide _______________ 22. magnesium bromide _____________ 21.Al2O3 22.MgBr2

23. sodium chlorate ________________ 24. iron II chloride ________________ 23.NaClO3 24.FeCl2

25. hydrogen gas __________________ 26. silver chromate ________________ 25. H2 26. Ag2CrO4

27. zinc bicarbonate _______________ 28. barium oxide ________________ 27.Zn(HCO3)2 28.BaO

29. aluminum nitrate ______________ 30. diphosphorus pentoxide __________ 29.Al(NO3)3 30.P2O5

31. aluminum hydroxide ___________ 32. chromium III oxide _____________ 31.Al(OH)3 32.Cr2O3

33. lithium phosphate ________________ 34. ice ________________________ 33. Li3PO4 34. H2O

35. nitrogen dioxide _________________ 36. iron III oxide _________________ 35. NO2 36. Fe2O3

37. sodium peroxide ________________ 38. copper II oxide ________________ 37.Na2O3 38.CuO2

39. liquid nitrogen _______________ 40. lead II acetate _________________ 39.N2 40.Pb(C2H3O2)2

41. lead IV fluoride ________________ 42. ferrous bromide ________________ 41. PbF4 42. FeBr2

43. carbonic acid _________________ 44. silver bisulfite ________________ 43.H2CO3 44.AgHSO3

45. cupric hydroxide ________________ 46. nitric acid __________________ 45.Cu(OH)2 46.HNO3

47. mercury II bromide _______________ 48. stannic sulfide ________________ 47. HgBr2 48. SnS2

49. hydrofluoric acid _______________ 50. potassium phosphate _____________ 49. HF 50. K3PO4

51. iodine tribromide _______________ 52. phosphorus pentafluoride __________ 51. IBr3 52. PF5

Contributors

• David Wilner (http://chemtutor.com/heurist.htm)

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